Chemical Equilibrium

A. Iron- Silver Equilibrium

The first part of the experimentation focuses in the iron-silver system. Silver nitrate (AgNO3) was added to ferrous sulfate (FeSO4) shown in this equation: 2AgNO3(l) + FeSO4(l) Ag2SO4-(aq)+ Fe(NO3)2(aq)

The additon of AgNO3to FeSO4 equilibrium had an effect on the equlibria. Though no precipitate was formed after just mixing the solutions, a precipitate formed after centrifugation. This was done to achieve separation of silver precipitate to the supernate.

The silver precipitate is formed due to the unsolubility of AgNO3 which displaces the chemical equilibrium of FeSO4 towards the silver compound. This is an application of Le-Chatelier fundament which states that an equilibrium will attempt to shift in a direction that will counteract a stress that is placed on it.

The supernate was then tested for the presence of three ions namely, Fe2+, Fe3+, and Ag+. In the test for Fe2+, he supernate was added with K3Fe(CN)6 that resulted to net ionic equation: Fe2+(aq) + K3Fe(CN)6(aq) KFe[Fe(CN)6](s)

The precipitate which is KFe[Fe(CN)6](s) is a prussian blue precipitae which indicates the presence of Fe2+ ions. Next, for Fe3+, KCNS was added in the supernate shown in this net ionic equation: Fe3+(aq) + SCN-(aq) Fe(SCN)2+(aq)

A blood red solution was achieved which indicates the presence of Fe3+ ions. Lastly, for the Ag+ test, the HCl was added to the supernate shown in the ionic equation below: Ag+(aq) + Cl-(aq) AgCl(s)

A white precipitate formed which indicates the presence ogf Ag+ ions.

Since the reactants and the products are both present at the same time, the system has reached its equilibrium. This means though products are produced, the reverse reaction still proceeds. For that it could be concluded that the range of Keq is between 10-2 and 102, significant amounts of both product and reactant will remain on the reaction and these substances, depending on the particular chemical reaction, can be easily be detected via qualitative test.

B. Copper-Ammonia Equilibrium

The second part of the experimentation focuses on the copper(II)-ammonia system. Ammonia (NH3) was added to cupric sulfate as shown in this chemical equation: (with limited NH3) CuSO4(aq) + NH3(aq) Cu(OH)+(s)

Cu(OH), a light blue precipitate formed upon adding smaller amount of ammonia. The precipitate formed because the concentration of the reactants were not equal thus no equilibrium attained. But with excess amount of ammonia added, as shown in this equation: (with excess NH3) CuSO4(aq) + 4NH3(aq) Cu(NH3)42+(aq)

The precipitate would dissolve due to the formation of soluble complex compound. The reaction was forced to go on completion which resulted to a deep blue complex.

When HCl was added, the production of NH4+ is triggered. The decolorization of the solution is due to the system at equilibrium given by the reaction below: NH3(aq)+H+(aq)⇆NH4+(aq)

The addition of HCl removes the available NH3 that will lead to the formation of the precipitate, copper (II) hydroxide and the complex ion, tetraamminecopper(II). Therefore, by the addition of HCl, the general chemical equilibrium is favored in the direction of the reverse reaction.

The amount of HCl added to shift the equilibrium to the reverse direction is 1/3 less than the amount of NH3 needed to complete the equilibrium (HCl- 10 drops, NH3-15 drops). This indicates that it is easier for HCl to shift the direction of the equilibrium in reverse than for NH3 which needed more amount to complete the reaction and attain equilibrium.

C. Chromate-Dichromate Equilibrium

The third part of the experimentation focuses on chromate-dichromate system. To start off, four wells were filled up with potasium chromate (K2CrO4) and potasium dichromate (K2Cr2O7) separately occupying two wells each. The original color of chromate was yellow while dichromate was orange.

This is a dynamic equilibrium and it is sensitive to the acidity or basicity of the solution. Shifting the equilibrium with pH changes is a classic example of Le Chatelier’s principle at work. It states that if a chemical dynamic equilibrium is disturbed by changing the conditions (concentration, temperature, volume or pressure), the position of equilibrium moves to counteract the imposed change. So if more reactant is added, the equilibrium shifts to the right in order to consume that extra reactant, which results in more product; also if the product is removed from the system, the equilibrium shifts to the right completely increasing the yield.

Upon adding H2SO4 (acid) in chromate, its color changed to orange as shown in the equation below: 2CrO42-(aq) + 2H+ (aq) ⇆ Cr2O72-(aq) + H2O(l)

As the sulfuric acid is added to the chromate solution, the yellow color turns to orange. Increasing the hydrogen ion concentration is shifting the equilibrium to the left in accordance with Le Chatelier’s principle, where we expect the reaction to try remove some of the H+ we have added by reacting with the CrO42-, and yielding more Cr2O72- which we observe as color change.

For the dichromate, NaOH (base) was added and its color changed to yellow as shown in the equation below: Cr2O72-(aq) + 2OH-(aq) ⇆ 2CrO42-(aq) + H2O(l)

When sodium hydroxide is added to the dichromate solution, the orange color turns back to yellow, hydroxide ions react with hydrogen ions forming water, driving the equilibrium to the right (OH- removes H+ ions by neutralizing them and the system acts to counteract the change) and further shifting the color. We can observe that the addition of hydroxide ions promotes the conversion of dichromate to chromate.

Sulfuric acid and sodium hydroxide was used to produce acidic and basic conditions for the solution because they dissociate completely in aqueous solution, enough to provide hydronium and hydroxide ions for the system at equilibrium.

NaOH was also added in chromate but there were no visible change. Therefore it could be concluded that chromate is stable under basic conditons. H2SO4 was also added in dichromate but there were no visible changes. Therefore it could be concluded that dichromate solution is stable under acidic conditions.

D. Iron-Thiocyanate Equilibrium

The fourth part of the experimentation focuses on the iron (III) chloride- thiocyanate system. Fe3+ and SCN- reacted as shown in the equation below: Fe3+(aq) + SCN-(aq) ⇆ Fe(NCS)2+ (aq)

The interaction of these compounds is an equilibrium producing a pale orange solution.

Using the principle of Le Chatelier, we could explain to which direction the equilibrium will shift when a certain compound is added as a reactant. According to this principle, a stress that involves an increase in the reactant concentration or decrease in product concentration will cause the reaction to move to the right, and a decrease in the reactant concentration or an increase in a product concentration will cause it to move to the left.

Adding either Fe3+ or SCN- to the reaction caused the color of the solutin to darken because, in order to relieve stress of added reactant concentration, more Fe(NCS)2+ was formed to use up all the reactants to still be able to establish equilibrium.

It would be a different case if NaCl was added to the reaction. The color of the solution would be lighter because NaCl removed some free SCN- as the complex, rather than destroying the complex directly and once again equilibrium was restored by dissociation of some Fe(NCS)2+, resulting to disappearance of some solution color. Therefore the direction of the equilibrium shifted to the left.

E. Cobalt-Cobalt Chloride Equilibrium

The fifth part of the experimentation focuses on the cobalt (II) ions system. A balanced equation of the reaction between Co2+ and Cl- can be seen below: Co2+ (aq) + 4Cl-(aq) (aq) ⇆ CoCl42-

If an aqueous solution contains both cobalt(II) and chloride ions, the blue ion CoCl42- forms, in equilibrium with the pink Co2+(aq) ion.

On test tube no. 1, without the addition of HCl, the equilibrium lies to the far left and the solution is pink. After adding HCl, the solution changed its color into blue because the equilibrium shifted to the right. The reaction is known to be sensitive to temperature. That is why at room temperature the solution in test tube no. 2 was pink, meaning the equlibrium shifted to the left. And upon heating the solution, its color changed to blue violet, meaning the equilibrium shifted to the right.

It can be concluded that the reaction was endothermic. Heating the solution makes the solution turn blue, favoring the formation of the blue complex, CoCl42-. The forward reaction is favored upon an increase in temperature. This observation suggests that the reaction is endothermic.

CONCLUSION AND RECOMMENDATION(S)

After conducting five sytems involving equilibrium reactions, the priciples that govern this phenomenon of equilibria was proven to be true, giving live applications of these principles. In an equilibrium reaction, many factors are considered to be able to state the presence of equilibrium and its shifting.

The concept and application of equilibrium is a helpful tool since the prediction and the results of equilibrium shifts of reaction, the favored conditions, and Keq range were able to explain the extent of a given reaction in terms of how much or how less products like medicine, polymer, and fuel can be obtained from a particular reaction mixture. It could also tell how can a reaction adjust conditions to produce more.

It would also be interesting to know how different species of compounds can have the same effect on the equilibrium of the reaction, or how a single species could cause a large effect on equilibrium reactions.